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8-1 Types of Chemical Bonds
))(1031.2( 2119
rQQ
nmJE
p330
Coulomb’s law
The energy of interaction between a pair of ionscan be calculated using Coulomb’s law:
where E has units joules, r is the distancebetween the ion centers in nanometers, andQ1 and Q2 are the numerical ion charges.
For example, the distance between the centers ofthe Na+ and Cl- ions is 0.276 nm, and the ionicenergy pair of ions is
Jnm
nmJE 1919 1037.8]276.0
)1)(1([)1031.2(
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5
Questions to Consider
What is meant by the term “chemical
bond?”
Why do atoms bond with each other to
form molecules?
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Figure 8.1(b) Energy profiles as a function between the
hydrogen atoms. As the atoms approach each other (right side
of graph), the energy decreases until the distances reaches
0.074 nm and then begins to increase again due to repulsions.
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Bondlength
Key ideas in bonding
Ionic Bonding: Electrons are transferred
Covalent Bonding: Electrons are shared
equally
What about intermediate cases?
Polar covalent bond:
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H F
FH F
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Chemical bond
What is meant by the term “chemical bond?”
Why do atoms bond with each other to form
molecules?
How do atoms bond with each other to form
molecules?
React 1 p333
10
8-2 Electronegativityp333
Expected H-X bonding energy= ½ (H-H bond energy + X-X bond energy)
△ = (H-X)act - (H-X)exp
The greater is the difference in the electronegativities of the atoms,the greater is the ionic compound and the greater is the value of △.
Linus Pauling(1901-1995)
If X has a greater electronegativity than H, the shares electron(s)will tend to be closer to the X atom. The molecule will be polar,with charge distribution.
Electronegativity: the ability of an atom in amolecule to attract shared electrons to itself.
Figure 8.2 The effect of an electric field on hydrogen fluoride molecules
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13
The Pauling electronegativity values
Figure 8-3
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The Pauling electronegativity values. Electronegativity generally
increases across a period and decreases down a group.
14
The general trend for electronegativity
What is the general trend for
electronegativity across rows and down
columns on the periodic table?
Explain the trend.
React 2 p334
15
Ex 8.1 Relative Bond PolaritiesP335
Order the following bonds according to polarity:
H–H, O–H, Cl–H, S–H, and F–H.
21
Ex 8.2 Bond Polarity and DipoleMoment
P337
For each of the following molecules, show the direction
of the bond polarities and indicate which ones have a
dipole moment: HCl, Cl2, SO3(a planar molecule with the
oxygen atoms spaced evenly around the central sulfur
atom), CH4 [trtrahedral(see Table 8.2) with the carbon
atom at the center], and H2S (V-shaped with the sulfur
atom at the point).
Choose an alkali metal, an alkaline metal, a noble gas,
and a halogen so that they constitute an isoelectronic
series when the metals and halogen are written as
their most stable ions.
What is the electron configuration for each species?
Determine the number of electrons for each species.
Determine the number of protons for each species.
Rank the species according to increasing radius.
Rank the species according to increasing ionization
energy.
React 3
29
What we can “read”from theperiodic table:
Trends for Atomic size Ion radius Ionization energy Electronegativity
Electron configurations Predicting formulas for ionic
compounds Ranking polarity of bonds
Ex 8.4 Relative Lon Size IIP342
Choose the largest ion in each of the followinggroups.
a.Li+, Na+, K+, Rb+, Cs+
b.Ba2+, Cs+, I-, Te2-
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Ex 8.3 Relative Lon Size IP342
Arrange the ions Se2-, Br-, Rb+, and Sr2+ in order of
decreasing size.
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Formation of an ionic solid
1. Sublimation of the solid metal
•M(s) M(g) [endothermic] (For Li(s) is +161 kJ.)
2. Ionization of the metal atoms
•M(g) M+(g) + e[endothermic] (For Li(g) is +520 kJ)
3. Dissociation of the nonmetal
•1/2X2(g) X(g) [endothermic] (For F is +½ (154 kJ)
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35
4. Formation of Xions in the gas phase:X(g) + e X(g) [exothermic] (For F- is -328 kJ/mole)
5. Formation of the solid MX:
M+(g) + X(g) MX(s) [quite exothermic]
(Corresponding to the lattice energy for LiF, which is -1047
kJ./mole)
Formation of an ionic solid (continued)
The relationship between the ionic character
of a covalent bond and the electronegativity
difference of the bonded atoms.
Figure 8.13
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42
8-7 The Covalent Chemical Bond:A Model
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Models
Models are attempts to explain how natureoperates on the microscopic level based onexperiences in the macroscopic world.
The Localized Electron Bonding Model
43
Fundamental Properties of Models
1. A model does not equal reality.
2. Models are oversimplifications, and aretherefore often wrong.
3. Models become more complicated as they
age.
4. We must understand the underlyingassumptions in a model so that we don’tmisuse it.
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Ex 8.5 △H from Bond Energies
Using the bond energies listed in Table 8.4, calculate △H
for the reaction of methane with chlorine and fluorine to
give Freon-12(CF2Cl2).
)(2)(2)()(2)(2)( 22224 gHClgHFgClCFgFgClgCH
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49
8-9 The Localized ElectronBonding Model
A molecule is composed of atoms that are
bound together by sharing pairs of electrons
using the atomic orbitals of the bound atoms.
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50
Localized Electron Model
1. Description of valence electron
arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).
3. Description of atomic orbital types used to
share electrons or hold long pairs.
51
8-10 Lewis Structure
Shows how valence electrons are arranged
among atoms in a molecule.
Reflects central idea that stability of a
compound relates to noble gas electron
configuration.
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52
Lewis Structures
1. Sum the valence electrons.
2. Place bonding electrons between pairs of
atoms.
3. Atoms usually have noble gas
configurations.
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Ex 8.6 Writing Lewis Structures P357
Give the Lewis structure for each of the following.
a. HF, b. N2, c. NH3, d. CH4, e. CF4, f. NO+
Ex 8.7 Lewis Structures for MoleculesThat Violate the Octet Rule I
P360
Write the Lewis structure for PCl5.
57
Ex 8.8 Lewis Structures for MoleculesThat Violate the Octet Rule II
P361
Write the Lewis structure for each molecule or ion.
a. ClF3 b. XeO3 c. RnCl2 d. BeCl2 e. ICl4-
Ex 8.9 Resonance StructuresP363
Describe the electron arrangement in the nitrite
anion (NO2-) using the localized electron model.
Ex 8.10 Formal ChargesP366
Give possible Lewis structures for XeO3 , an explosive
compound of xenon. Which Lewis structure or structures
are most appropriate according to the formal charges?
63
8-13 Molecular Structure:The VSEPR Model
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VSEPR Model
The structure around a given atom is
determined principally by minimizing
electron pair repulsions.
65
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put pairs as far apart as possible.
3. Determine positions of atoms from the way
electron pairs are shared.
4. Determine the name of molecular structure
from positions of the atoms.
P369Ex 8.11 Prediction of MolecularStructure I
Describe the molecular structure of the water molecule.
The Lewis structure for water is
There are four pairs of electrons: two bonding pairs and twononbonding pairs. To minimize repulsions, these bestarrangement in a tetrahedral array, as shown in Fig. 8.17.
Solution
Figure 8.17
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Ex 8.12 Prediction of Molecular Structure II
When phosphorus reacts with excess chlorine gas, the
compound phosphorus pentachloride (PCl5) is formed. In
the gaseous and liquid states, this substance consists of
PCl5 molecules, but in the solid state it consists of a 1 : 1
mixture of PCl4+ and PCl6
- ions. Predict the geometric
structures of PCl5, PCl4+, and PCl6
-.
SolutionThe Lewis structure for PCl5 is shown. Five pairs of electrons aroundthe phosphorous atom require a trigonal bipyramidal arrangement(see Table 8.6).
The Lewis structure for the PCl4+ ions (5+4(7) -1 = 32 valence
electrons) is shown. There are four pairs of electrons surroundingthe phosphorus atom in the PCl4
+ ion, which requires a tetrahedralarrangement of the pairs.
The Lewis structure for PCl6- (5 + 6(7) + 1 = 48 valence electrons)
is shown. Since each electron pair is shared with a chlorine atom, anoctahedral PCl6
- anion is predicted.
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73
P373
Ex 8.13 Prediction of MolecularStructure III
Because the noble gases have filled s and p valence orbitals,they were not expected to be chemically reactive. In fact,for many years these elements were called insert gasesbecause of this supposed inability to form any compounds.However. In the early 1960s several compounds of krypton,xenon, and radon were synthesized. For example, a team atthe Argonne National Laboratory produced the stablecolorless compound xenon tetrafluoride (XeF4). Predict itsstructure and whether it has a dipole moment.
Solution
The Lewis structure for XeF4 is
The xenon atom in this molecule is surrounded bysix pairs of electrons, which means an octahedralarrangement.
The arrangement in Fig. 8.20(b) is preferred, and themolecular structure is predicted to be square planar. Thereis an octahedral arrangement of electron pairs, but theatoms form a square planar structure. Although each Xe-Fbond is polar, their structure causes the polarities to cancel.Thus XeF4 has no dipole moment as shown in the margin.
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P376Ex 8.14 Structures of Molecules withMultiple Bonds
Predict the molecular structure of the sulfur dioxide
molecule. Is this molecule expected to have a dipole moment?Solution
We must determine the Lewis structure for the SO2molecule, which has 18 valence electrons. The expectedresonance structures are
The structure of the SO2 molecule expected to be V-
shaped, with a 120-degree bond angle. The molecule has a
dipole moment as shown: